Chapter 7 - Gases, Liquids, And Solids
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Slide 7-2 Kinetic Molecular Theory of matter
Five statements of the kinetic molecular theory of matter:•1. matter is ultimately composed of tiny particles (atoms, molecules or ions) that have
definite and characteristic sizes that do not change
•
2. the particles are in constant random motion and therefore possesses kinetic energy.•3. the particles interact with one another through attractions and repulsions and
therefore possesses potential energy.
•
4. The kinetic energy of the particle increases as the temperature is increased•5. The particles in a system transfer energy to each other through elastic collisions.•In elastic collision total kinetic energy is conserved. In a inelastic collision some of the
energies are lost during collision. In real world every collision is inelastic.
•
Slide 7-3 Kinetic Molecular Theory of matter
Kinetic energy is energy that matter possess because the motion of the particle. It can be
considered as disruptive force that tends to make the particles of a system increasingly independent of one another.
•
Potential energy is stored energy that matter possesses as a result of its position,
condition or composition. Potential energy of attraction can be considered as cohesive force that tends to cause order and stability among the particles of the system.
•
Electrostatic attraction is an attraction or repulsion that occurs between charged particle.
It is a form of potential energy
•
Slide 7-4 Solids
When liquids are cooled, their molecules come so close together and attractive forces
between them become so strong that random motion stops and a solid is formed.
•
A solid is a physical state characterized by a dominance of potential energy (cohesive
forces) over kinetic energy (disruptive forces).
•
Characteristic properties of solids•1. Definite volume and definite shape: the strong, cohesive forces hold the particles in
essentially fixed positions, resulting in definite volume and definite shape
•
Slide 7-5 Solids
2. high Density: the constituent particles of solids are located as close together as
possible (touching each other). Therefore a given volume contains large number of particles , resulting in high density.
•
3. Small compressibility: Because there is very little space between particles, increased
pressure cannot push the particles any closer together; therefore it has little effect on solid’s volume.
•
4. Very small thermal expansion: An increased temperature increase the kinetic energy
(disruptive forces), thereby causing more vibrational motion of the particles. Each occupies a slightly larger volume, and the result is a slight expansion of the solid. The
strong cohesive forces prevent this effect from becoming very large.
•
CH 7 - Gases, Liquids, and SolidsWednesday, November 14, 20074:38 PM
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Slide 7-6 Liquids
As pressure increases in a real gas, its molecules come closer and closer with the result
that attractions between molecules become important.
•
When distances decrease so that almost all molecules touch or almost touch, a gas
condenses to a liquid.
•
A liquid is the physical state characterized by potential energy (cohesive forces) and
kinetic energy (disruptive forces) of about the same magnitude.
•
Characteristic properties of liquids:•1. Definite volume and indefinite shape: The attractive forces are strong enough to
restrict particles to movement within a definite volume. They are not strong enough, however to prevent the particles from moving over each other in a random manner that
is limited only by the container walls. Thus liquids have no definite shape except that
they maintain a horizontal upper surface in container that are not completely filled
•
Slide - 7-7 Liquids
2. High Density: The particles in a liquid are not widely separated; they are still touching
one another, therefore a large number of particles in a given volume- a high density
•
3. Small compressibility: Because the particles in a liquid are still touching each other,
there is a very little empty space. Therefore an increase in pressure cannot squeeze the particles much closer together.
•
4. Small thermal expansion: Most of the particle movement in a liquid is vibrational
because a particle can move only a short distance before they collide with a neighbor. The increased particle velocity that accompanies a temperature increase results only in
increased vibrational amplitudes. The net effect is an increase in the effective volume a
particle occupies, which causes a slight volume increase in the liquid.
•
Slide 7-8 Gases
A gas is the physical state characterized by a complete dominance of kinetic energy
(disruptive forces) over potential energy (cohesive forces).
•
Properties of gases:•1. Indefinite volume and indefinite shape: the attractive forces between particles have
been overcome by kinetic energy, and the particles are free to travel in all directions. Therefore , gas particles completely fill their container, and the shape of the gas is that of
the container.
•
2. Low density: The particles of a gas are widely separated. There are relatively few
particles in a given volume (compared with liquids and solids), which means little mass per volume.
•
Slide 7-9 Gases
3. Large Compressibility: particles in a gas are widely separated; essentially a gas is
mostly empty space. When pressure is applied the particles are easily pushed closer together, decreasing the amount of empty space and the volume of the gas.
•
4. Moderate thermal expansion: An increase in temperature means an increase in particle
velocity. The increased kinetic energy of the particles enables them to push back whatever barrier is confining them into a given volume, and the volume increases. Size of
the particle is not changed during expansion or compression, they move further apart or
closer. It is the distance between two particles that changes.
•
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Slide 7-10 Gases
Gas law is the generalization that describes in mathematical terms the relationships
among the amount, pressure, temperature and volume of a gas
•
most commonly measured in millimeters of mercury (mm Hg), atmospheres (atm),
and torr.
•
1 atm = 760 mm Hg = 760 torr
= 101,325 pascals = 28.96 in. Hg
=14.7psi (lb/in2+)•pressure is measured using a barometer (next screen).•
Gas pressure: the pressure per unit area exerted against a surface.•
Slide 7-11 Gas Pressure
Figure 6.2 A mercury barometer.•
Slide 7-12 Gas Pressure
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Slide 7-12 Gas Pressure
Figure 6.3 A manometer.•
Slide 7-13 Gas Laws
PV = constant or P1V1 = P2V 2
Boyle’s law: for a fixed mass of gas at a constant temperature, the volume is inversely
proportional to the pressure.
•
V
T
V1
T1
V2
T2
= a constant or =
Charles’s Law: the volume of a fixed mass of gas at a constant pressure is directly
proportional to the temperature in kelvins (K).
•
Slide 7-14 Gas Law
Gay-Lussac’s Law: for a fixed mass of gas at constant volume, the pressure is directly
proportional to the temperature in kelvins (K).
•
P
T
P1
T1
P2
T2
= a constant or =
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T
V 2
T2
V 1
T1
P2
T2
P1
T1
P
V
Name Expression Constant
Boyle's law
Charles's Law
Gay-Lussac's law
P1V 1 = P2V2
=
=
in summary:•
Slide 7-15 Gas Laws
P1 V1
T1
P2 V2
T2
PV
T== a constant or
Boyle’s law, Charles’s law and Gay-Lussac’s law can be combined into one law called the
combined gas law.
•
Slide 7-16 Gas Laws
P1 = 2.00 atm V1 = 3.00 LInitial:
Final: P2 =10.15 atm V2 = ?
Problem: a gas occupies 3.00 L at 2.00 atm. Calculate its volume when the pressure is
10.15 atm at the same temperature.
•
because the temperature is constant T1 = T2•
P1 V1T2
T1 P2
V2 = =(2.00 atm)(3.00 L)
10.15 atm= 0.591 L
Slide 7-17 Gas Laws
The actual temperature and pressure at which we compare two or more gases does
not matter.
•
For convenience in making comparisons, chemists have selected one pressure as a
standard pressure, and one temperature as a standard pressure.
•
The standard temperature and pressure (STP) selected are 0°C (273 K) and 1 atm
pressure.
•
Avogadro’s law: equal volumes of gas at the same temperature and pressure contain the
same numbers of molecules.
•
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Slide 7-18 Gas Laws
All gases at STP or any other combination of pressure and temperature contain the same
number of molecules in a given volume. But how many is that?
•
One mole contains 6.022 x 1023 formula units; what volume of gas at STP contains this
many molecules?
•
This volume has been measured and found to be 22.4 L.•Thus, one mole of any gas at STP occupies 22.4 L•
Slide 7-19 Ideal Gas Law
Avogadro’s law allows us to write a gas law that is valid not only for any P, V, and T but
also for any mass of gas.
•
PV = nRTP = pressure of the gas in atmospheres (atm)
V = volume of the gas in liters (L)
n = moles of the gas (mol)
T = temperature in kelvins (K)
R = ideal gas constant (a constant for all gases)
Ideal gas law:•
Slide 7-20 Ideal Gas Law
We find the value of R by using the fact that 1.00 mol of any gas at STP occupies 22.4 L•
PVR =
nT=
(1.00 atm)(22.4 L)
(1.00 mol)(273 K)= 0.0821
L• atmmol• K
Problem: 1.00 mol of CH4 gas occupies 20.0 L at 1.00 atm. What is the temperature
of the gas in kelvins?
•
Solution: solve the ideal gas law for T and plug in the given values:•
PVnR
T = = 244 K(1.00 atm)(20.0 L)
(1.00 mol)(0.0821 L• atm• mol-1• K
-1)
=
Slide 7-21 Gas Laws
Dalton’s law of partial pressures: the total pressure, PT, of a mixture of gases is the sum
of the partial pressures of each individual gas:
•
PT = P1 + P2 + P3 + . . .
Problem: to a tank containing N2 at 2.0 atm and O2 at 1.0 atm we add an unknown
quantity of CO2 until the total pressure in the tank is 4.6 atm. What is the partial pressure of CO2?
•
4.6 atm 2.0 atm 1.0 atm 1.6 atm
Totalpressure
Partial pressure
of N 2
Partial pressure
of O2
Partial pressureof CO2
+ +=
Solution:•
Slide 7-22 Gases
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Slide 7-22 Gases
Ideal gas: the six assumptions of the KMT give us an idealized picture of the particles of a
gas and their interactions with one another.
•
their atoms or molecules do occupy some volume.•there are forces of attraction between their atoms or molecules.•
Real gases•
at pressures above 1 to 2 atm and temperatures well above their boiling points,
real gases behave in much the same way as predicted by the KMT.
•In reality, no gases are ideal•
Slide 7-23 Changes of States
A change of state is a process in which a substance is transformed from one physical
state to another physical state. They are normally accomplished by heating or cooling.
•
There are six possible changes of states: •Freezing: liquid to solid state•Melting: solid to liquid state•Evaporation: liquid to gaseous state•Condensation: gaseous to liquid state•Sublimation: solid to gaseous state•Deposition: gaseous to solid state•Endothermic Change: change of state in which heat energy is absorbed. Ex: melting,
evaporation, sublimation
•
Exothermic change: change of state in which heat energy is given off. Ex: freezing,
condensation, deposition
•
Slide 7-24 Intermolecular Forces
At or near STP, the forces of attraction between molecules of most gases are so
small that they can be ignored.
•
When T decreases or P increases or both, the forces of attraction become important
to the point that they cause condensation (gases to liquids) and ultimately solidification (liquids to solids).
•
In order to understand the properties of liquids and solids, we must look at the
nature of these intermolecular forces of attraction.
•
The strength of attractive forces between molecules determines whether any sample of
matter is a gas, liquid, or solid.
•
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Slide 7-25 Intermolecular Forces
Their origin is electrostatic; that is, the attraction between positive and negative
dipoles.
•
The strengths of covalent bonds are shown for comparison.•
We discuss three types of intermolecular forces•
Slide 7-26 London Dispersion Forces
London dispersion forces are the attraction between temporary induced dipoles.•
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Slide 7-27 London Dispersion Forces
London dispersion forces exist between all atoms and molecules.•
They are the only forces of attraction between atoms and nonpolar molecules.•
They range in strength from 0.01 to 2 kcal/mol depending on mass, size, and shape of
the interacting molecules.
•
In general, their strength increases as the mass and number of electrons in a molecule
increases.
•
Even though these forces are very weak, they contribute significantly to the attractive
forces between large molecules because they act over large surface areas.
•
Slide 7-28 Dipole-Dipole Interactions
CH3 CH2 CH2 CH3 CH3 -C- CH3
O-
+Butane
(bp 0.5°C)
Acetone
(bp 58°C)
consider butane and acetone, compounds of similar molecular weight•
Butane is a nonpolar molecule; the only interactions between butane molecules are
London forces.
•
Acetone is a polar molecule; its molecules are held together in the liquid state by
dipole-dipole interactions.
•
Dipole-dipole interactions; the electrostatic attraction between positive and negative
dipoles.
•
Slide 7-29 Hydrogen Bonds
Hydrogen bond: a noncovalent force of attraction between the partial positive charge on
a hydrogen bonded to an atom of high electronegativity, most commonly O or N, and the partial negative charge on a nearby O or N.
•
H O
H
H
O
H
- +
hydrogenbond
hydrogen
bond
- +
(a) (b) (c)
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Slide 7-30 Hydrogen Bonds
The strength of hydrogen bonds ranges from 2 to 10 kcal/mol.•
that in liquid water is approximately 5.0 kcal/mol•
By comparison, the strength of an O-H covalent bond in a water molecule is 119
kcal/mol.
•
Nonetheless, hydrogen bonding in liquid water has an important effect on the physical
properties of water.
•
The relatively high boiling point of water is due to hydrogen bonding between water
molecules; extra energy is required to separate a water molecule from its neighbors.
•
Hydrogen bonds are not restricted to water; they form whenever there is are O-H or N-H
groups.
•
Slide 7-31 Evaporation
In a liquid there is a distribution of kinetic energies (KE) among its molecules.•
Some have high KE and move rapidly; others have low KE and move more slowly.•
If a molecule at the surface is moving slowly (has a low KE), it cannot escape from
the liquid because of the attractions of neighboring molecules.
•
If, however, it is moving more rapidly (has a higher KE) and moving upward, it can
escape the liquid and enter the gaseous space above it.
•
An important property of liquids is that they evaporate:•
Slide 7-32 Evaporation
If the container is open, this process continues until all molecules escape.•
If the container is closed, molecules remain in the air space above the liquid.•
At equilibrium, molecules continue to escape from the liquid while an equal number are
recaptured by it.
•
The partial pressure of the vapor in equilibrium with the liquid is called the vapor
pressure of the liquid.
•
Vapor pressure is a function of temperature. Vapor pressure increases with temperature
until it equals the atmospheric pressure.
•
Evaporation is the process by which molecules escape from the liquid phase to the gas
phase. Evaporation is surface phenomena.
•
A vapor is a gas that exists at a temperature and pressure at which it ordinarily would be
thought of as a liquid or solid
•
Slide 7-33 Vapor Pressure of Liquids
Vapor pressure is the pressure exerted by a vapor above a liquid when the liquid and the
vapor are in equilibrium with each other.
•
A volatile substance is a substance that readily evaporates at room temperature because
of a huh vapor pressure. Gasoline is a substance whose components are highly volatile.
•
Boiling is a form of evaporation where conversion from the liquid state to the vapor state
occurs within the body of the liquid through bubble formation.
•
Boiling point is the temperature at which the vapor pressure of a liquid becomes equal to
the external (atmospheric) pressure exerted on the liquid
•
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Slide 7-34 Boiling Point
Normal boiling point: the temperature at which the vapor pressure of a liquid equals the
atmospheric pressure.
•
Slide 7-35 Conditions that affect Boiling Point
Boiling point of a liquid can be increased by increasing the external pressure. This
principle is used in pressure cooker cooking. Normally at high altitude cooking is slow because water boils at low temperature due to low pr. But in pressure cooker water boils
above 1000C at elevated pr. So the cooking is faster than normal.
•
Liquids that have high normal boiling points or undergo undesirable reaction at a higher
temperature can be made to boil at low temperatures by reducing the external pressure. This principle is used in numerous food products like juice preparation where water is
boiled off at reduced pr. Thus concentrating the juice without heating it at high
temperature.
•
Slide 7-36 Boiling Point
CHCl3
CH3 CH2 OH
H2 O
CH3 COOH
120
46
18
60
62
78
100
118
Chloroform
Ethanol
Water
Acetic acid
NameMolecular
Formula
Molecular
Weight
(amu)
Boiling
Point
(°C)
CH3 CH2 CH2 CH2 CH2 CH3 86 69Hexane
Table 6.3•
Boiling points of covalent compounds depend primarily on two factors: (1) the nature and
strength of intermolecular forces and (2) molecular size and shape.
•
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Slide 7-37 Boiling Points
Consider CH4 (MW 16, bp -164°C) and H2O (MW 18, bp 100°C). The difference in
boiling points between them is due to the greater strength of hydrogen bonding in water compared with the much weaker London dispersion forces in methane.
•
Consider methane, CH4 (MW 16, bp -164°C), and hexane C6H14 (MW 86, bp 69°C).
Because of its larger surface area, London dispersion forces are stronger between hexane molecules than between methane molecules.
•
Intermolecular forces•
Slide 7-38 Boiling Points
when molecules are similar in every way except shape, the strength of London forces
determines boiling point
•
CH3 CH2CH2CH2 CH3CH3 -C-CH3
CH3
CH3
Pentane(bp 36.2°C)
2,2-Dimethylpropane(bp 9.5°C)
Both are C5H12 and have the same molecular weight.•
2,2-dimethylpropane is roughly spherical while pentane is a linear molecule.•
Pentane has the higher boiling point because it has the larger surface area and
stronger London dispersion forces between its molecules.
•
Molecular shape•
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